1. Reactions in Aqueous Media

2. General Properties of Solutions
A solution is a homogeneous mixture of two or more substances
The substance present in the greater amount is the solvent
The substance present in the smaller amount is the solute
A solution may be a gas (e.g., air), solid (e.g., an alloy) or liquid (e.g., seawater)
Aqueous solutions are of greatest interest to us

3. Electrolytes and Non-electrolytes
An electrolyte is a substance which, in water, forms an electrically conductive solution
A non-electrolyte is a substance which, in water, forms a solution which is not electrically conductive
An electrically conducting solution must contain ions so that the electric current can pass from one electrode to the other
Pure water contains very few ions:

4. Electrolytes and Non-electrolytes
An electrolyte dissociates to form ions in solution
a strong electrolyte dissociates 100% into ions when dissolved in water, for example,
The polarity of the water molecules helps stabilize ions
Hydration is the process by which the water molecules are arranged in a particular way around the ions

5. Electrolytes and Non-electrolytes
Ionic compounds are strong electrolytes
A strong acid is a strong electrolyte because it completely dissociates in aqueous solution
A weak acid does not dissociate completely in aqueous solution and is therefore a weak electrolyte
Because the dissociation of a weak electrolyte is incomplete, for the same molarity, a solution of a weak electrolyte is not as good of a conductive solution as that of a strong electrolyte

6. Precipitation Reactions
A precipitation reaction is characterized by the formation of an insoluble product, i.e., a precipitate
A precipitation reaction usually involves ionic compounds
You can predict whether a precipitate will form if you know the solubility of the ionic compounds
Three categories of solubility for compounds: soluble, slightly soluble, and insoluble

7. Solubility 1. All alkali metal compounds are soluble.
2. All ammonium compounds are soluble.
3. All nitrates, chlorates, and perchlorates are soluble.
4. Most hydroxides are insoluble, except for alkali metal hydroxides and barium hydroxide. Calcium hydroxide is slightly soluble.
5. Most chlorides, bromides, and iodides are soluble, except those that form Ag+, Hg22+, or Pb2+ ions.
6. All carbonates, phosphates, and sulfides are insoluble, except those bound to alkali metals or ammonium ions.
7. Most sulfates are soluble. Calcium sulfate and silver sulfate are slightly soluble. Barium sulfate, mercury(II) sulfate and lead sulfate are insoluble.

8. Solubility
Example: Decide whether these compounds are soluble, slightly soluble or insoluble:
Ag2SO4
CaCO3
Na3PO4
Solutions:
Slightly soluble according to rule 7
Insoluble according to rule 6
Soluble according to rule 1

9. Ionic Equations
There are several ways to write a precipitation reaction
For example, is a molecular equation
A molecular equation is useful because it tells us what solutions we must combine to form the precipitate
However, the molecular equation is misleading because the BaCl2(aq), Na2SO4(aq), and NaCl(aq) all dissociate in water and do not remain intact

10. Ionic Equations
An ionic equation is closer to reality because it represents the ionic compounds in the form of free ions
In our example, Na+ and Cl- are spectator ions (they do not participate in the reaction in any way)
The net ionic equation neglects the spectator ions and indicates that the species directly involved in the reaction

11. Ionic Equations
Example: Decide what precipitate will form in the following reaction and give the net ionic equation for this reaction:
Solution: According to rule 4, aluminum hydroxide is insoluble. The molecular equation is The ionic equation is The net ionic equation is

12. Acid-Base Reactions Arrhenius’ definitions of acids and bases:
An acid ionizes in water to give H+ ions
A base ionizes in water to give OH- ions
Bronsted’s definitions of acids and bases:
A Bronsted acid is a proton donor
A Bronsted base is a proton acceptor

13. Acid-Base Reactions
In water, a proton, H+(aq), will actually associate with a molecule of water to form the hydronium ion, H3O+(aq)
We will use H+(aq) and H3O+(aq) interchangeably
H+(aq) is more convenient
H3O+(aq) is what we really have in solution
The ionization of an acid can therefore be written in two ways:

14. Acid-Base Reactions
For Bronsted acids, there are acids that can release one, two, or three protons, making them monoacids, diacids, an triacids respectively
N.B. the double arrow is used to indicate that the acid is weak and the dissociation is incomplete

15. Acid-Base Reactions
Ionic compounds that posses OH- anions are Bronsted bases since the OH- ion can accept a proton
Ammonia (NH3) is a weak Bronsted base
N.B. Ammonium hydroxide, NH4OH, does not exist: in reality, it is an aqueous solution of ammonia

16. Acid-Base Reactions
Example: Decide whether the following species are Bronsted acids or bases: (a) HBr, (b) NO2-, (c) HCO3-, (d) SO42-, (e) HI
Solution: (a) acid: (b) base: (c) acid: base: HCO3-(aq) is amphoteric (acid and base) (d) base: (e) acid:

17. Neutralization Reactions
A neutralization reaction, a reaction between an acid and a base, has the general form: acid + base salt + water, where the salt is made up of a cation other than H+ and an anion other than OH- or O2-
ex.; Molecular equation: Ionic equation: Net ionic equation:

18. Oxidation-Reduction Reactions
An oxidation-reduction, or redox, reaction involves the transfer of electrons
Each redox reaction has a reducing agent and an oxidizing agent
A reducing agent donates electrons
An oxidizing agent receives electrons or, in other words,
An oxidizing agent captures electrons
A reducing agent loses electrons

19. Oxidation-Reduction Reactions
ex.; The reaction 2 Ca(s) + O2(g) CaO(s) is a redox reaction
Ca(s) is the reducing agent
O2(g) is the oxidizing agent
it is useful to separate a redox reaction into two half- reactions
Oxidation reaction (loss of electrons): 2 Ca Ca e-
Reduction reaction (gain of electrons): O e O2-
the number of electrons involved in each half reaction must be identical

20. Oxidation-Reduction Reactions
The reaction of a metal with an acid is a redox reaction:
metal + acid salt + H2(g)
ex.;
The sum of the half reactions is the net ionic equation
Oxidation:
Reduction:

21. Oxidation States
For ionic compounds, the concepts of oxidation and reduction are obvious as there is a complete transfer of electrons
For covalent compounds, there is only a partial transfer of electrons
We introduce the concept of oxidation states
The oxidation state of an atom in a molecule corresponds to the charge it would have if there was a complete transfer of electrons from the less electronegative atom to the more electronegative atom

22. Oxidation States
ex.;
The numbers that appear above the elements are the oxidation states
An oxidation is the increase in the oxidation state of an element
A reduction is the decrease in the oxidation state of an element
A redox reaction is thus any reaction where there are changes in oxidation states

23. Oxidation States (Rules)
1. For a single element, each atom has an oxidation state on zero
For monatomic ions, the oxidation state is equal to the charge on the ion. All alkali metals have a +1 oxidation state and all alkaline earth metals have a +2 oxidation state, regardless of the compound. Aluminum has an oxidation state of +3 in its compounds.
3. The oxidation state of oxygen in most compounds is -2; however, in hydrogen peroxide (H2O2) and peroxide ions (O22-), its oxidation state is -1.

24. Oxidation States (Rules)
The oxidation state of hydrogen is +1, except when bound to a metal in a binary compound. In this case, its oxidation state is -1.
Fluorine has an oxidation state of -1 in its compounds. Other halogens have negative oxidation states when they appear as a halide ion in a compound. However, when they combine with oxygen, they have positive oxidation states.
6. In a neutral molecule, the sum of oxidation states of all atoms must be zero. In a polyatomic ion, the sum of oxidation states of all elements must be equal to the net charge of the ion.

25. Oxidation States
Example: Determine the oxidation states of each element in the following compounds and ions: (a) Li2O, (b) Na2O2, (c) C6H12O6, (d) BaH2, (e) OCl-.
Solution: (a) According to rule 2, Li is According to rule 6 (or 3), O is (b) According to rule 2, Na is According to rule 6 (or 3), O is (c) According to rule 3, O is According to rule 4, H is According to rule 6, C is (d) According to rule 2, Ba is According to rule 6 (or 4), H is (e) According to rule 3, O is According to rule 6, Cl is +1

26. Oxidation States

27. Oxidation States Metallic elements only have positive oxidation states
Non-metals can have negative or positive oxidation states
The representative atoms (groups IA to VIIA) cannot have an oxidation state that is higher than the number of their group
Transition metals usually have several possible oxidation states

28. The Concentration of Solutions
The concentration of a solution is the amount of solute in a given amount of solution
The molarity is the number of moles of solute contained per unit volume of solution (in liters)
Molarity is an intensive property

29. The Concentration of Solutions
Example: What is the molar concentration of an aqueous solution containing 1.77 g of ethanol (C2H5OH) in a volume of 85.0 mL?
Solution:

30. The Dilution of Solutions
A dilution is the process that reduces the concentration of a solution
Over the course of a dilution, the number of moles of a solute does not change
i.e., moles of solute before dilution = moles of solute after dilution
i.e., Cinitial Vinitial = Cfinal Vfinal

31. Gravimetric analysis
Gravimetric analysis is an analytical method based on the measurement of the mass of a precipitate
For the analysis to be precise, the precipitation reaction must be complete, i.e., the precipitate that forms must be highly insoluble
For example, to know the quantity of Cl- in solution, we would add excess AgNO3 to the solution to produce the AgCl(s) precipitate
If we weigh the amount of AgCl(s) produced, we can calculate the amount of Cl-(aq) that was dissolved

32. Gravimetric analysis
Example: A g sample of an ionic compound containing bromine ions (Br-) is dissolved in water. To this, we add an excess amount of AgNO3. If the mass of the AgBr(s) produced is g, what is the weight percent of Br in the initial compound?
Solution: We formed g of AgBr,
So we had mol of Br- in solution or ( mol)(79.90g/mol) = g . The weight percentage of Br in the initial compound is:

33. Acid-Base Titrations moles of standard = moles of unknown
In a titration , a solution of known concentration (the standard solution) is added gradually to a solution of unknown concentration until the chemical reaction between the two solutions is complete, i.e., it has reached the equivalence point
determined with the help of an indicator
If one knows the exact volumes of standard and unknown solution, one can determine the concentration of the unknown solution
moles of standard = moles of unknown

34. Acid-Base Titrations moles of CH3COOH = moles of NaOH
Example: What is the concentration of an acid solution (CH3COOH) if its volume is 50.0 mL and we need 27.2 mL of a standard solution of NaOH (0.245 mol/L) to neutralize the acid solution?
Solution:
moles of CH3COOH = moles of NaOH

35. Acid-Base Titrations moles of OH- = moles of H+
Example: How many milliliters of a 1.28 mol/L H2SO4 solution are needed to neutralize 60.2 mL of a mol/L KOH solution?
Solution: N.B. H2SO4 is a diacid, so 2 moles of KOH are stoichiometrically equivalent to1 mole of H2SO4: we must take into account that the acid has twice as many H+ ions to donate
moles of OH- = moles of H+